EXCHANGE 


Addition  Compound  Formation  In 
Aquedus  Solutions; 

Hydrates  at  the  Boiling-Point. 


Q 


DISSERTATION 

SUBMITTED    IN    PARTIAL    FULFILLMENT    OF    THE    RE- 
QUIREMENTS FOR  THE  DEGREE  OF  DOCTOR  OF 
PHILOSOPHY  IN   THE  FACULTY  OF  PURE 
SCIENCE  OF  COLUMBIA  UNIVERSITY 


BY 


HORATIO  WALES,  Jr.,  A.  B.,  M.  A. 

r 

31T>3 


Tri-County  Press,  Polo,  Illinois 
1921 


Addition  Compound  Formation  In 
Aqueous  Solutions; 

Hydrates  at  the  Boiling-Point. 


— o— 


DISSERTATION 

SUBMITTED    IN    PARTIAL    FULFILLMENT    OF    THE    RE- 
QUIREMENTS FOR  THE  DEGREE  OF  DOCTOR  OF 
PHILOSOPHY   IN   THE  FACULTY  OF  PURE 
SCIENCE  OF  COLUMBIA  UNIVERSITY 


BY 

HORATIO  WALES,  Jr.,  A.  B.,  M.  A. 


Tri-County  Press,  Polo,  Illinois 
1921 


ACKNOWLEDGEMENT 

To  Professor  James  Kendall,  at  whose  suggestion  this  prob- 
lem was  undertaken,  the  author  wishes  to  express  his  sincere 
gratitude  for  help  and  advice  during  the  progress  of  the  work. 

The  author  also  desires  to  extend  his  thanks  to  Professor 
Thos.  B.  Freas,  for  the  many  helpful  suggestions  made  during 
Professor  Kendall's  absence. 


*-'- 

Addition  Compound  Formation  In  Aqueous 
Solutions;   Hydrates  at  tne  Boiling  Point. 


For  many  years  a  large  number  of  chemists  have  believed 
that  combination  often  occurs  in  a  solution  between  the  solvent 
and  the  solute.  The  methods  used  to  show  the  existence  of  com- 
pound formation  in  solution  are  many  and  varied.  The  absorp- 
tion spectra,1  specific  heat,2  viscosity,3  surface  tension,4  in  fact 
the  deviations  which  a  large  number  of  dissolved  substances 
show  from  nearly  every  law  of  physics  and  physical  chemistry 
have  been  used  to  uphold  the  solvate  theory-5 

This  investigation  was  undertaken  with  the  purpose  of  demon- 
strating the  existence  of  hydrates  by  the  boiling-point  method. 
Very  little  work  has  been  done  along  this  line,  four  chemists 
having  investigated  non-aqueous  solutions6  and  only  one 
aqueous  solutions.7  As  all  of  these  workers  have  attempted  to 
compute  the  amount  of  solvation  by  the  "molecular  rise  of  the 

iJones  and  Uhler,  Am.  Chem.  Jour.  37,  126,  207,  244,  (1907); 
Jones  and  Anderson,  Ibid,  41,  163,  276,  (1909);  Jones  and  Strong, 
Ibid,  45,  1,  (1911);  Moore,  Zeit.  Phys.  Chem.  55,  641,  (1906);  Lewis, 
Ibid,  56,  223,  (1906). 

aTimofejew,  Chem.  Cent.  2,  429,  (1905);  Bose,  Zeit.  Phys.  Chem. 
58,  585,  (1907);  Muller  and  Fuchs,  Compt.  Rend.  140,  1639,  (1905); 
Dupre  and  Page,  Proc.  Roy.  Soc.  20,  336,  (1872). 

sDunstan  and  Thole,  Trans.  Chem.  Soc.  95,  1556,  (1909);  Smith 
and  Menzies,  Jour.  Am.  Chem.  Soc.  31,  1191,  (1909);  Gaillaud  and 
Bingham,  Am.  Chem.  Jour.  35,  195,  (1906). 

4Linebarger,  Jour.  Am.  Chem.  Soc.  22,  5,  (1900);  Whatmough, 
Zeit.  Phys.  Chem.  39,  129,  (1901);  Grunmach  Ann.  Phys.  (4)  9,  1261, 
(1902);  Morgan,  Jour.  Am.  Chem.  Soc.  37,  1416,  (1915);  Ibid.  39, 
2261,  2275,  (1917). 

sAbout  twenty  methods  in  all  have  been  used.  For  a  complete 
list  of  investigators  and  methods  used  see  Bauer,  Ahrens  Sammlung,  8, 
466,  (1903);  Washburn,  Tech.  Quart.  21,  363,  (1908);  Dhar,  Zeit. 
Elektrochem,  20,  57,  (1914). 

eJones  and  Getman,  Am.  Chem.  Jour.  32,  338,  (1904);  Werner, 
Zeit.  Anorg.  Chem.  15,  1,  (1895);  Biltz,  Zeit.  Phys.  Chem.  40,  185, 
(1902);  Walden,  Ibid,  55,  321,  (1906). 

TJohnston,  Trans.  Roy.  Soc.  Edin.  45,  193,  (1906). 


.  : 

* 


,    .     i  «•  d-    i    *     * 


boiling-point,"  their  results,  as  we  shall  show  presently,  are  of 
little  value. 

The  boiling-point  method  has  also  been  used  to  determine 
molecular  weights,1  vapor  pressure  lowering,2  heats  of  dilution,3 
and  ionization.4  With  one  exception5  all  investigators  have 
used  the  van't  Hoff  equation6  to  compute  their  results. 

In  a  number  of  communications  recently  appearing,7  addition 
compounds  of  organic  acids  with  organic  substances  containing 
oxygen  have  been  described  and  compared.  In  a  later  paper8  the 
results  of  this  series  of  investigations  were  summed  up  and  cer- 
tain laws  derived,  which  were  then  applied  to  aqueous  solutions 
of  a  series  of  acids  and  their  validity  tested  by  means  of  freezing- 
point  curves.  As  the  results  of  these  investigations  are  of  great 
importance  and  have  a  direct  application  to  this  work,  it  seems 
advisable  to  repeat  them  here. 

When  an  organic  substance  containing  oxygen  is  added  to 
an  acid  of  the  type  HX,  the  existence  of  an  addition  compound 
in  the  liquid  mixture  can  usually  be  demonstrated.  In  every 
case  the  formation  of  addition  compounds  becomes  more  evident 
both  in  the  number  and  the  stability  of  the  compounds  isolated, 
the  more  divergent  the  acidic  strengths  of  the  two  components. 
With  substances  of  similar  acidic  strengths  no  evidence  of  com- 
pound formation  has  been  found. 

In  the  system  water-acid  the  conditions  were  shown  to  be 
exactly  analogous  to  those  for  acids  in  pairs.  With  water  func- 


iq.  v.  Landolt-Bornstein,  Tabellen. 

aRaoult,  Compt.  Rend.  103,  1125,  (1887);  Dieterici,  Wied.  Ann 
62,  616,  (1897);  Smits,  Amsterdam,  2,  469,  (1910). 

ajuttner  Zeit.  Phys.  Chem.  38,  112,   (1901). 

4Baroni,  Gazz.  Chim.  Ital.  23,  [1],  263;  [2],  249,  (1893);  Sal- 
vador!, Ibid,  26,  [1],  237,  (1896);  Schlamp,  Zeit.  Phys.  Chem.  14, 
273,  (1894). 

5Juttner,  loc.  cit. 

eThe  usual  form  of  the  van't  Hoff  equation  for  the  elevation  of 
the  boiling-point  is  DT=Kc,  where  K  is  a  constant  determined  by  the 
solvent  and  c  the  weight  concentration  of  the  solvent. 

7Kendall,  Jour.  Am.  Chem.  Soc.  36,  1222,  1722,  (1914);  38,  1309, 
(1916);  Kendall  and  Gibbons,  Ibid,  37,  149,  (1915);  Kendall  and 
Carpenter,  Ibid,  36,  2498,  (1914);  Kendall  and  Booge,  Ibid,  38,  1712, 
(1916). 

sKendall,  Booge  and  Andrews,  Jour.  Am.  Chem.  Soc.  39,  2303, 
(1917). 


tioning  as  an  extremely  weak  acid,  no  evidence  could  be  observed 
of  hydrate  formation  with  exceedingly  weak  acids.  As  the 
strength  of  the  acid  was  increased,  a  regular  increase  could  be 
seen  in  the  extent  of  combination  in  the  liquid  state  ;  the  stronger 
the  acid,  the  greater  the  amount  of  hydration.  This  need  not  be 
taken  as  the  only  proof  that  hydration  increases  with  the  acidic 
strength.  The  literature  is  filled  with  instances  of  hydrate  for- 
mation. A  few  examples  may  well  be  cited  here.1 

r-(CHOHCOOH)a;  H*0 
(COOHCH2)2COHCOOH;  H20 

(COOH)a;2ILO 

C6H5SO3H;  H2O 
CeIL(S03H)2;  2.5HX) 
;  3ILO 


H3As04;  0.511,0 

IKPO*;  H.O 

H3P04;  0.1H2O;  O.SIM) 

HC1;H20;  2H.O;  3H.O 

HC1O4;  ILO;  2H*0;  2.5H2Q  ;  SILO;  3-5ILO 

ILSO<;  HX);  2H*O;  4ILO 

These  examples  clearly  illustrate  that  hydrate  formation  is 
the  rule.  In  every  case  the  number  and  complexity  of  the  hydrates 
increases  with  the  strength  of  the  acid.  These  acids  must  exist 
in  solutions  largely  combined  with  the  solvent  or  their  hydrates 
would  not  separate  out  as  the  solid  phase.  No  hydrates  of  very 
weak  acids,  either  organic  or  inorganic,  have  been  isolated.  If 
these  acids  as  a  class  are  combined  with  the  solvent  to  any  ex- 
tent, some  of  their  hydrates  would  surely  be  known.  However, 
the  failure  to  isolate  a  hydrate  cannot  be  taken  as  proof  of  its 
non-existence.  Compounds  may  be  present  in  the  solution  and 
yet  not  be  stable  either  at  the  temperatures  used  for  evaporation 
or  freezing. 

The  existence  of  hydrates  in  solution  has  already  been  indi- 
cated for  a  series  of  acids  by  means  of  freezing-point  curves.2 
On  account  of  the  slight  solubility  very  few  acids  are  available 


KDrganic  acids  from  Beilstein;   inorganic  from  Landolt-Bornstein, 
Tabellen,   (1912),  p469-472. 

2Kendall,  Booge,  and  Andrews,  loc.  cit. 


6 

for  freezing-point  determinations  at  high  concentrations.  At  the 
boiling-point,  however,  the  increased  solubility  allows  a  much 
wider  choice  of  material.  For  this  reason,  and  also  since  few 
solid  hydrates  are  stable  above  ordinary  temperatures,  it  was 
thought  advisable  to  see  if  the  relationships  which  have  been 
demonstrated  for  solutions  at  low  temperatures  are  also  valid  at 
the  boiling-point.  We  would  expect,  however,  since  the  most 
complex  hydrates  which  have  been  isolated  were  formed  at  the 
lowest  temperatures,  that  compound  formation  would  be  less 
extensive  at  the  boiling-point  than  at  any  lower  temperature. 

THEORETICAL  BASIS  OF  THE  METHOD 

It  is  a  well  known  fact  that  the  simple  van't  Hoff  equation 
for  the  elevation  of  the  boiling-point  can  be  used  only  for  infin- 
itely dilute  solutions-1  The  true  expression  for  the  boiling 
point  of  a  solution  of  a  non-volatile  substance,  valid  at  all  con- 
centrations, is  readily  derived  from  a  thermodynamical  investi- 
gation of  the  problem. 

If  to  a  solution  of  a  non-volatile  solute  at  its  boiling  point  T 
we  add  dN  moles  of  solute,  thus  increasing  its  mol  fraction  by  dy 
and  decreasing  that  of  the  solvent  by  dx,  the  vapor  pressure  will 
vary  as  the  mol  fraction  of  the  solvent  in  the  solution,2  and  will 
be  given  by  the  expression: 

(1)  p=kx 

differ entiating,  we  obtain 

(2)  dp=kdx 

and  dividing  2  by  1  we  find  that  the  decrease  in  vapor  pressure 
of  the  solvent  is  given  by  the  expression; 

(3)  dp/p=dx/x 

From  the  Clapeyron-Clausius  equation  we  find  that  for  small 
temperature  intervals 

lEven  in  very  dilute  solutions  this  formula  cannot  be  considered 
as  giving  consistent  results.  Many  investigators  have  attempted  to  de- 
termine molecular  weights  from  boiling-point  measurements,  even 
though  no  two  compounds  nor  two  concentrations  of  the  same  com- 
pound give  identical  values  for  the  "molecular  elevation  of  the  boil- 
ing-point." For  example,  different  investigators  have  found  values 
of  this  "constant"  for  aqueous  solutions  of  KC1  ranging  from  463  to 
621.  (Johnston,  loc.  cit.  p.  195.) 

sRaoult,  loc.  cit.  The  mol  fraction  of  the  solvent  is  denned  as  the 
ratio  of  the  mols  of  solvent  to  the  total  mols  in  the  solution. 


—  RTMp 
(4)  dT  £ 

L.    p 

Substituting  the  value  of  dp/p  in  this  equation  we  obtain:1 

—  RT2dx 

(5)  dT=      — 

L       x 

where  x  is  the  mol  fraction  of  the  solvent  in  the  solution  ;  T  the 
absolute  temperature  of  the  boiling-point  of  the  pure  solvent;  L 
the  molar  heat  of  vaporization  of  the  pure  solvent  (537.5x18.016)  ; 
and  R  the  gas  constant,  1.9885  calories. 

In  order  to  integrate  this  equation  it  is  necessary  first  to 
express  L  as  a  function  of  AT  which  can  only  be  done  approx- 
imately. We  are  fortunate,  however,  in  having  another  solution 
of  the  problem.  If  L  in  a  given  case  can  be  regarded  as  practically 
constant  over  the  range  covered  by  AT,  the  integral  becomes: 

—  0.4343L     AT 

(6)  log,»x=—  ^  — 

where  TB  is  the  boiling-point  of  the  solution.  This  equation  may 
also  be  written: 

—  L    AT 

(7)  loge  (l-y)_-  -- 

where  y  is  the  mol  fraction  of  solute.  Expanding  log*  (1  —  y)  by 
means  of  Taylors  Theorem  we  obtain: 

(8)  loge  (1—  y)=loge  l—  y/l!+y2/2!—  y3/3!+  ...... 

When  the  solution  becomes  very  dilute  y  becomes  very  small  and 
we  may  disregard  all  the   higher   powers.   Equation   8   then   be- 
comes : 

(9)  loge  (l—  y)  =—  y 

The  temperature  interval  will  also  become  so  small  that  TTB  is 
nearly  equal  to  T2.  Equation  7  may  then  be  written  : 

(10)  y=  t    AT 


which  is  the  usual  form  of  the  van't  Hoff  equation.2  This  equation 
represents  the  limit  approached  as  the  solution  becomes  more  and 
more  dilute.  By  no  means  can  the  use  of  a  formula  which  only 


ic.  f.  Washburn.  Principles  of  Physical  Chemistry,  p.  164. 
2c.  f.  Ikeda,  Jour.  Col.  Sc.  Tokyo,  29,  no.  10,  p.  21,  (1908). 


8 

holds  for  infinitely  dilute  solutions  be  justified  for  computing 
the  hydrates  formed  in  solutions  of  finite  concentrations. 

From  equation  6  we  see  that  if  the  boiling-point  of  a  solu- 
tion is  abnormally  raised  it  will  mean  that  the  value  which  was 
assigned  to  x  is  smaller  than  that  calculated  from  the  weights  of 
the  two  components  taken.  If  we  have  made  the  corrections  re- 
quired by  the  ionization  and  the  change  in  L,  the  only  explana- 
tion we  are  able  to  give  at  the  present  time  is  that  some  of  the 
solvent  has  been  removed  by  combination  with  the  solute  and 
thus  caused  its  mol  fraction  to  be  lowered.  The  actual  number  of 
solvent  molecules  combining  with  each  molecule  of  solute  can  be 
found  by  determining  what  value  must  be  given  to  x  to  make 
the  observed  values  of  AT  fall  on  the  ideal  curve.  The  compar- 
ison of  the  variation  in  the  boiling-point  elevation  of  solutions  of 
a  series  of  compounds  of  varying  strengths  will,  therefore,  afford 
a  means  of  indicating  their  relative  degree  of  hydration  in  solu- 
tion. 

APPLICATION  OF  THE  THEORY  TO  THIS  WORK 

The  best  methods  that  have  been  devised  for  determining 
the  rise  in  boiling-point  of  a  solvent,  caused  by  the  presence  of 
a  dissolved  substance,  are  not  free  from  many  objections,  both 
in  theory  and  method.  Some  of  these  apparently  cannot  be  elim- 
inated, while  others  can  be  overcome  to  a  greater  or  less  degree. 
With  an  electrolyte  as  the  solute  our  systems  are  far  from  ideal. 
The  many  disturbing  factors  influencing  x  and  T  must  be  taken 
into  account.  For  the  purpose  of  discussion  these  will  be  divided 
into  three  classes ;  those  due  to  the  solvent,  those  due  to  the 
solute,  and  those  due  to  the  method. 

FACTORS  DUE  TO  THE  SOLVENT 

Since  water  is  an  associated  liquid,  the  use  of  the  value 
H2O=18.016  in  the  calculations  of  the  mol  fractions  will  lead  to 
values  for  x  which  are  much  too  high.  It  would  seem  at  first 
sight  as  though  there  is  no  solution  to  this  difficulty.  Fortunately, 
however,  the  variation  in  x  with  the  molecular  weight  assumed 
for  water  is  largely  counterbalanced  on  the  right  hand  side  of 
the  equation  by  the  fact  that  L,  the  molar  heat  of  vaporization 
is  similarly  affected.  The  actual  differences  in  y,  the  mol  frac- 


tion  of  solute,  and  AT  as  calculated  for  given  concentrations  of 
solute  from  equation  7  for  two  of  the  molecular  forms  of  water 
are  given  in  the  following  table: 

Mols  solute  in    (a)  Water=H>O  (b)  Water=  (H*0).  Difference 

lOOOgms.  water    y          AT  y              AT              in    AT 

0.5607                 0.01      0.288  0.0199      0.288                 0.000 

1.1328                 0.02      0.578  0.0392      0.572                 0.006 

2.3127                 0.04      1.170  0.0769       1.155                 0.015 

3.5429                 0.06      1.776  0.1132       1.724                 0.042 

It  is  apparent  from  these  figures  that  the  variation  in  the  ideal 
curve  with  the  association  factor  of  water  is  beyond  the  ordinary 
experimental  errors.  At  the  boiling-point  water  is  mostly  dihy- 
drol  with  a  small  amount  of  monohydrol.1  Since  the  association 
factor  of  water  also  varies  with  the  concentration  of  the  solute, 
the  water  molecule  becoming  less  complex  in  the  more  concen- 
trated solutions,  it  is  evident  that  basing  the  theoretical  curve 
on  the  assumption  that  water=H2O  would  render  all  subsequent 
calculations,  such  as  the  estimation  of  the  water  of  hydration, 
merely  approximations,  except  in  very  dilute  solutions.2 

In  our  work  the  relative  positions  of  the  curves  are  all  we 
can  hope  to  establish  and  the  mol  fractions  can  all  be  expressed 
on  the  basis  of  water  =H2O,  since  this  will  give  a  constant  error 
throughout  the  range  of  concentrations  studied. 

In  the  integration  of  the  above  equation  we  assumed  that 
the  heat  of  vaporization  of  water  remains  constant  over  the 
range  of  temperatures  studied  (about  three  degrees).  The  fact 
that  the  heat  of  vaporization  will  change  slightly  with  the  tem- 
perature leads  to  a  complication  in  the  integrated  equation. 

We  may  express  L  as  a  function  of  the  temperature  by  the 
equation : 

(11)  L=L0+  (Cp— C)T 


iThe  constitution  of  water.     Tr,  Farad.  Soc.  6,  71,  (1910). 

'Previous  workers  have  shown  little  hesitation  on  this  point. 
Johnston  (loc.  cit.  p.  873)  calculates  the  hydration  when  AT  is  53° — 
and  by  the  van't  Hoff  equation. 


10 

where  L0  is  the  heat  of  vaporization  at  absolute  zero.1 
Substituting  the  numerical  values  for  L0  (which  is  assumed  to  be 
constant)  and  the  molar  specific  heats  at  100°C,  the  fundamental 
equation  becomes; 

dx  13574—  10.43T 

(12)  _  =  --  dt 
x  1.9885T2 

Integrating  this  equation  we  obtain  ; 


(13)  logex=5.246  loge- 

T        T          TB 

Prom  these  two  equations  we  then  obtain  the  following  theoreti- 
cal values  for  the  two  highest  concentrations  studied. 

From  6  From  13 

AT  x  x 

1.471  0.95000  0.95005 

1.776  0.94000  0.94006 

From  these  values  it  can  be  seen  that  the  effect  of  correcting  for 
the  change  in  L  with  the  temperature  is  infinitesimal. 

The  value  of  L  also  changes  slightly  with  the  pressure.  The 
error  due  to  this  cause  was  entirely  eliminated  in  this  work  by 
always  using  a  pressure  of  760  millimeters  of  mercury. 

ERRORS  DUE  TO  SOLUTE 

The  total  number  of  molecules  in  a  solution  of  an  electrolyte 
will  be  increased  by  ionization.  If  we  know  the  ionization  at  any 
given  concentration  we  can  at  once  calculate  the  mol  fraction 
of  the  solute.  The  exact  estimation  of  ionization  in  solutions  of 
the  high  concentrations  examined  here  is  subject  to  many  errors. 
The  conductance  ratio  in  these  cases  cannot  indicate  the  true 
degree  of  dissociation  and  a  large  viscosity  correction  is  neces- 
sary2 which  at  best  will  be  only  approximate.  Moreover  for  strict 
accuracy,  each  solution  requires  conductivity  and  viscosity  data 


iKendall  (Meddel.  fran  K.  Vets.-Akad.  Nobel,  Band  2,  no.  36, 
(1913)  )  has  found  that  the  heat  of  vaporization  for  one  gram  of  wat- 
er is  more  nearly  expressed  by  the  equation  Q=83.06  (Tc-T)%,  where 
TC  is  the  critical  temperature  for  water,  370<>C.  On  account  of  the 
very  complicated  integral  this  equation  gives,  it  could  not  be  used. 

2Kraus  and  Bray,  Jour.  Am.  Chem.  Soc.  35,  1315,  (1913.) 


11 

for  its   own  particular  temperature   of  ebullition,   whereas,   all 
conductivity  measurements  were  made  at  100°C. 

At  high  concentrations  these  errors  become  large,  but  since 
the  ionization  in  these  cases/vis  extremely  small,  their  actual  ef- 
fect is  minimized,  and  their  total  influence  on  x  is  inconsiderable. 
For  weak  acids  the  ionization  oc  may  be  calculated  directly 
from  the  equation  i1 

ccc2 


where  c  is  the  weight  concentration  of  the  solute  and  k,  the  ioni- 
zation constant.  For  stronger  acids,  such  as  phosphoric  and 
oxalic,  oc  is  more  nearly  represented  by  the  equation:2 

(1—  oc) 


1  —  cc  oc 

where  k  and  k7  are  constants  depending  on  the  electrolyte. 

Another  possible  disturbing  factor  is  the  association  of  the 
solute  in  solution.  This  would  tend  to  place  the  value  of  the  mol 
fraction  of  the  solute  below  that  calculated  on  the  assumption 
of  simple  molecules.  Except  in  a  few  rare  cases  association  has 
only  been  found  to  exist  in  very  concentrated  solutions  and  here 
it  is  known  to  decrease  with  a  rise  in  temperature.3  As  the  mol 
fraction  of  the  solute  in  this  work  was  always  small,  extensive 
association  is  not  to  be  expected  and  errors  due  to  this  cause  may 
be  neglected  at  the  boiling  temperatures. 

In  the  integration  of  the  fundamental  equation  it  was  as- 
sumed that  L  is  independent  of  the  concentration.  This  is  true 
only  if  the  heat  of  dilution  is  zero  throughout  the  entire  range. 
As  this  is  not  the  case  the  corrections  necessary  in  the  integrated 
equation  must  be  taken  into  account. 

To  evaporate  one  gram-molecule  of  water  at  100°  C,  L  cal- 

iThis  equation  has  been  shown  to  hold  exactly  for  concentrations 
up  to  weight  normal.  Kendall,  Jour.  Am.  Chem.  Soc.  36,  1083  (1914). 

sKendall,  loc.  cit. 

3Peddle  and  Turner  (Jour.  Chem.  Soc.  99,  690,  (1911)  )  from 
investigations  on  the  boiling  and  freezing-points  of  solutions  of  a  large 
number  of  organic  substances  in  water,  have  found  that  the  only  class 
of  substances  which  show  no  evidence  of  association  is  the  aliphatic 
amines.  Marked  association  was  shown  only  by  the  aromatic  com- 
pounds. 


12 

ories  are  necessary.  If,  however,  to  one  mole  of  water  we  add  n 
moles  of  some  acid,  AQ  calories  will  be  given  off  or  absorbed. 
The  work  to  separate  the  solvent  from  this  solution  will  then 
become  equal  to  L+AQ  calories,  where  AQ  is  either  positive  or 
negative. 

Sufficient  data  is  not  available  at  the  present  time  to  calcu- 
late this  correction.1  However,  by  using  a  series  of  compounds 
which  have  heats  of  solution  of  the  same  order  of  magnitude,  we 
will  introduce  a  constant  error  in  our  work  and  the  curves  may 
be  used  to  show  relative  hydration,  although  no  comparison  can 
be  made  between  the  lowest  curve  and  the  theoretical  curve. 

In  view  of  the  above  errors  it  becomes  apparent  that  the  lim- 
itations of  the  method  are  of  such  a  magnitude  that,  although 
with  proper  choice  of  materials  we  can  show  the  relative  hydra- 
tion of  a  series  of  compounds,  it  would  be  utterly  absurd  to  at- 
tempt to  calculate  the  actual  amount  of  combination  between  the 
solvent  and  the  solute. 

ERRORS  DUE  TO  THE  METHOD 

The  boiling-point  method  is  not  capable  of  that  degree  of  ac- 
curacy to  which  the  freezing-point  method  has  been  developed. 
The  difference  between  the  temperature  of  the  liquid  and  the 
surrounding  objects  is  much  greater  than  in  the  freezing-point 
method  and  correspondingly  greater  precautions  are  necessary 
to  protect  the  solution  and  the  thermometer  from  changes  in  the 
temperature  of  external  objects. 

A  slight  change  in  pressure  will  cause  a  large  change  in  the 
boiling-point  of  a  solution  while  its  effect  on  the  freezing-point 
is  noticed  only  in  the  fourth  place  of  decimals.  Part  of  the  sol- 
vent is  continually  separating  as  vapor  and  is  returned  as  liquid 
at  a  temperature  much  lower  than  that  of  the  boiling  solution. 


lAbegg  (Zeit.  Phys.  Chem.  15,  247,  )1894)  )  has  found  that  "the 
corrections  for  the  heat  of  dilution  of  acetic,  formic,  and  tartaric  acids, 
for  depressions  up  to  15°  are  at  the  utmost  of  the  same  order  of  mag- 
nitude as  the  experimental  errors  in  the  freezing-point  determina- 
tions." When  we  consider  the  magnitudes  of  the  experimental  errors 
of  the  freezing  and  boiling  point  methods,  and  recall  that  an  elevation 
of  0.52°  in  the  boiling-point  is  equivalent  to  a  depression  of  1.86°  In 
the  freezing-point,  it  becomes  at  once  apparent  that  the  corrections 
due  to  the  heat  of  dilution  for  these  acids  at  the  boiling-point  cannot 
be  otherwise  than  negligible. 


13 

Although  the  loss  due  to  this  cause  cannot  be  determined  accur- 
ately, the  following  illustration  will  show  that  it  is  negligible. 

The  volume  of  the  vapoj;  above  the  liquid  was  about  lOOcc. 
Let  us  assume  that  the  concentration  of  the  solute  is  5% — calcu- 
lated on  the  basis  of  no  loss  of  vapor — and  that  the  weight  of 
water  lost  is  0.1  gm.,  a  value  which  is  much  too  high.  Now,  since 
we  started  with  thirty-five  grams  of  water  in  each  case,  we 
would  have  in  this  solution  0.1023  gram-molecules  of  solute.  Cor- 
recting for  the  loss  of  vapor  we  would  have  34.9  gms.  of  water 
acting  as  the  solvent,  which  would  make  the  concentration  of 
the  solute  equal  to  5.01%  molar.  We  have  purposely  chosen  a 
case  where  the  error  due  to  this  cause  would  be  at  a  maximum 
and  have  magnified  this  error.  As  this  concentration  represents 
the  upper  limit  of  the  curve  and  the  errors  in  all  other  cases  will 
be  smaller,  we  may  safely  neglect  this  factor  entirely. 

At  the  boiling-point  the  solution  must  be  in  equilibrium 
with  the  vapor  of  the  pure  solvent.  To  obtain  this  the  vapor  and 
the  solution  must  be  intimately  mixed,  regular  boiling  must  be 
insured,  and  the  heat  exchange  with  the  surrounding  objects 
must  be  at  a  minimum;  otherwise,  the  temperature  observed  will 
not  be  the  temperature  of  the  solution.  The  first  condition  is 
secured  when  the  solution  is  boiling  briskly  and  is  continually 
agitated  and  always  in  contact  with  the  bubbles  of  vapor  pass- 
ing upwards.  The  second  condition  is  accomplished  as  nearly  as 
possible  by  the  form  of  the  apparatus  used,  the  details  of  which 
will  be  given  later. 

METHOD 

Previous  investigators  have  found  that  when  water  is  used 
as  the  solvent,  the  boiling  point  results  were  not  satisfactory 
and  in  nearly  every  case  the  data  thus  obtained  has  been  used  to 
compute  molecular  weights.1  A  study  of  the  results  obtained 
shows  that  in  no  case  have  two  compounds  nor  two  concentra- 
tions of  the  same  compound  given  identical  values  for  the  "  mol- 
ecular rise  of  the  boiling-point.'*2 

iLandolt-Bornstein,  Tabellen,  loc.  cit. 

'Johnston  (loc.  cit.)  finds  that  the  results  of  different  investigat- 
ors vary  by  as  much  as  20%  and  thinks  that  this  may  he  due  to  errors 
in  the  determination  of  the  boiling-point  of  water.  The  magnitude  of 
the  error  makes  this  explanation  very  improbable. 


14 

It  would  seem  that  the  aqueous  solutions  have  presented  a 
special  difficulty  and  in  order  to  obtain  uniform  results,  nearly 
every  investigator  has  designed  a  special  form  of  apparatus,  a 
rule  to  which  we  are  no  exception. 

The  boiling-point  of  a  liquid  is  defined1  as  the  lowest  temper- 
ature at  which  a  liquid  can  remain  steadily  boiling  at  any  given 
pressure.  However,  it  has  been  found  by  several  investigators2 
that  for  the  attainment  of  exact  temperature  adjustment  it  is 
necessary  to  maintain  an  extremely  energetic  boiling.  Although 
this  may  lead  to  a  possible  source  of  error  (due  to  superheating), 
it  may  be  minimized  by  supplying  the  heat  to  the  solution  at  a 
constant  rate  during  the  boiling  (which  would  tend  to  keep  the 
amount  of  superheating  constant).  This  is  practically  impossible 
to  accomplish  if  we  are  to  use  the  usual  method  of  heating  with 
a  bunsen  burner. 

To  overcome  this  difficulty  somewhat,  we  used  as  our  source 
of  energy  a  cylindrical  resistence  heater,  into  which  the  tubes 
were  placed.  The  tubes  were  at  first  separated  from  the  heater 
by  an  air  space,  but  later  an  asbestos  cylinder  was  placed  inside 
the  heater  as  it  was  thought  that  the  sudden  changes  observed  in 
the  thermometer  readings  were  caused  by  heat  being  radiated 
into  the  bulb.  Whether  this  is  the  true  explanation  or  not,  no 
further  trouble  was  experienced  along  this  line. 

It  was  also  noticed  that  a  change  of  30  watts  (corresponding 
to  a  change  in  current  of  one  ampere)  in  the  power  supplied  to 
the  heater  would  cause  a  change  in  the  thermometer  readings  of 
about  O.°02;  the  temperature  remaining  constant  in  either  case. 
In  this  work  the  power  supplied  to  the  heater  was  kept  constant 
at  all  times,  thus  ensuring  a  constant  rate  of  boiling  in  all  the 
determinations. 

The  boiling-point  tubes  used  in  this  investigation  were  a 
modification  of  the  usual  Beckmann  type,  being  slightly  larger 
and  with  the  arms  placed  near  the  top  in  order  to  give  as  small 
a  vapor  space  as  possible.  Previous  investigators  have  used  gar- 
nets in  the  boiling  point  tubes  to  secure  a  uniform  ebullition. 
We  found  that  these  were  appreciably  soluble  in  the  boiling  solu- 

iNernst,  Theoretical  Chemistry. 

2Biltz,  Zeit.  Phys.  Chem.  40,  207,   (1902). 


tions  of  the  stronger  acids  used,  such  as  oxalic  and  phosphoric 
acids,  and  substituted  small  quartz  crystals  with  good  results. 
The  same  weight,  thirty  grams,  of  quartz  was  always  used  in 
order  to  keep  their  volume^  constant. 

The  boiling-point  tube  was  connected  to  a  large  water  man- 
ometer through  an  air  reservoir  (see  diagram),  by  means  of 
which  the  pressure  on  the  liquid  was  kept  constant  at  760  milli- 
meters of  mercury  throughout  each  determination.  The  heater  and 
the  boiling-point  tube  were  entirely  surrounded  with  asbestos 
boards  which  served  to  prevent  sudden  cooling  of  the  liquid 
by  air  currents  and  thus  prevent  sudden  changes  in  the  ther- 
mometer readings.  All  temperatures  were  measured  with  a 
Beckmann  thermometer  which  had  been  calibrated  by  the  U.  S. 
Bureau  of  Standards. 

To  prevent  the  liquid  which  had  been  cooled  by  the  conden- 
ser from  coming  in  contact  with  the  bulb  of  the  thermometer 
and  causing  sudden  changes  in  the  readings,  a  platinum  cylinder 
which  reached  well  above  the  surface  of  the  liquid,  was  placed 
around  the  bulb.1  This  also  served  to  reduce  any  radiation  of 
heat  from  the  bulb  of  the  thermometer  outwards. 

In  carrying  out  a  determination,  weighed  amounts  of  the 
acid  under  consideration  were  placed  in  the  boiling-point  tube 
and  thirty-five  grams  of  water  added.  Just  before  the  liquid 
came  to  a  boil,  its  level  was  noted  on  a  graduated  scale  placed 
behind  the  thermometer,  and  the  correction  for  the  immersed 
depth  of  the  thermometer  bulb  made.  The  density  of  the  solu- 
tion was  determined  with  sufficient  accuracy  for  this  from  the 
weights  of  water  and  acid  taken,  assuming  the  expansion  of  the 
solution  to  be  the  same  as  that  for  water.  The  barometric  pres- 
sure was  then  recorded  and  the  total  pressure  on  the  liquid  cor- 
rected to  760mm.  by  means  of  the  manometer.  Pressure  differ- 
ences of  less  than  one  millimeter  were  corrected  for  by  adding 
or  subtracting  0.° 00375  to  the  corrected  thermometer  readings 


iJones,  Am.  Chem.  Jour.  19,  584,  (1897). 


16 


17 

for  each  tenth  millimeter  of  mercury  differing  from  760.1  Ther- 
mometer readings  were  taken  at  intervals  of  thirty  seconds  over 
a  period  of  four  or  five  minutes  and  their  average  recorded.  By 
working  only  on  days  wfyen  the  barometer  was  fairly  constant, 
we  were  able  to  obtain  temperatures  which  did  not  vary  by  more 
than  O.°01  during  the  period  of  observation.  All  experiments 
were  repeated  at  least  twice  before  the  final  results  were  plotted. 
About  fifty  points  were  determined  for  the  curve  of  each  acid. 
The  method  of  computing  the  results  was  as  follows: 

Weight  acid  used    (Succinic) 2.3433  gms 

Weight    water  35.00       gms 

Gram  molecules  succinic  acid 0.0198 

Gram  molecules  water 1.9444 

Molar  per  cent  acid 1.01% 

Barometer  height  (corrected)   755.13      mm. 

Correction  due  to  immersed  depth  of 

thermometer2  2.65       mm. 

The  2.22  mm  required  to  bring  the  pressure  to  760mm  were 
added  by  means  of  the  manometer. 

Observed  boiling-point    (average  of  ten  readings 

taken  at  thirty  second  intervals) l.°115 

Stem  temperature  30°. 

Calibration    correction   .  — .  004 


1.  Ill 


iThis  correction  although  approximate  is  well  within  the  experi- 
mental error  for  the  pressure  differences  used.  Berkeley  and  Appleby, 
working  with  saturated  salt  solutions  found  that  the  effect  of  changes 
in  the  barometric  pressure  on  the  temperature  of  equilibrium  were 
greater  than  with  pure  water.  (Proc.  Roy.  Soc.  85A,  492,  (1911)  ) 
The  magnitude  of  the  coefficient  depends  upon  the  solubility  of  the 
salt  and  also  upon  the  temperature  coefficient  of  the  solubility.  Thus 
for  Na,SO4  which  is  only  moderately  soluble  and  does  not  vary  in  solu- 
bility with  the  temperature,  the  coefficient  is  0.°00052C  per  millimet- 
er of  mercury,  while  for  KNO3  which  has  both  a  high  solubility  and  a 
large  temperature  coefficient  of  solubility,  the  coefficient  is  0.°0217C 
per  millimeter,  or  about  forty  times  as  large.  This  plainly  shows  that 
the  method  of  actually  measuring  the  temperature  difference  between 
the  boiling-points  of  the  solution  and  pure  solvent,  regardless  of  the 
pressure,  cannot  possibly  give  correct  results  unless  the  pressure  is 
760  millimeters  on  both  liquids. 

sSmith  and  Menzies,  Jour.  Am.  Chem.  Soc.  32,  901,   (1910). 


18 

Reading  corrected  for  calibration 1.  Ill 

Setting  factor  (1°=1.0186)  .  020 


1.  131 
100°  corresponds  on  Beckmann  thermometer  to —  .  904 


0.  227 
Exposed  stem  correction .  018 


Boiling-point    elevation 0.  245 

CONDUCTIVITY  MEASUREMENTS 

The  conductivity  measurements  were  made  with  a  Preas 
cell  having  electrodes  approximately  one  centimeter  square  and 
one  centimeter  apart.  The  resistance  was  measured  with  a  Leeds 
Northrup  bridge  and  the  current  produced  with  a  Leeds  North- 
rup  Constant  High  Speed  Generator  which  gave  a  current  of  an 
approximately  sine  wave  form  and  a  frequency  of  1000  cycles. 
Temperature  regulation  was  obtained  by  placing  the  conductiv- 
ity cell  in  a  steam  bath,  allowing  about  one  half  hour  in  each 
case  for  the  solution  to  reach  the  temperature  of  the  steam.  Al- 
though the  temperature  of  boiling  water  may  vary  widely,  as 
has  been  mentioned,  the  vapor  above  it  will  always  have  the 
same  temperature;  this  temperature  being  determined  by  the 
barometric  pressure.  Except  in  extreme  cases  this  temperature 
will  not  vary  from  100° C  by  more  than  0.°2. 

In  computing  the  volume  occupied  by  the  solution  at  100°  it 
was  assumed  that  the  thermal  expansion  coefficients  of  the  solu- 
tions were  the  same  as  that  for  water.  Although  this  is  not 
strictly  true  it  is  within  the  experimental  error  for  the  small 
volumes  worked  with.  In  every  case  two  solutions  of  widely 
different  concentrations  were  prepared  and  the  dilutions  made 
on  both  in  order  to  minimize  any  errors  due  to  volume  measure- 
ments. 

The  cell  was  standardized  by  determining  the  conductivity 
of  a  solution  containing  ten  mini-equivalents  of  potassium  chlor- 
ide per  liter.  The  specific  conductivity  of  a  solution  of  this 
strength  at  100°  is  7.54xlO-3mhos.1  All  solutions  were  prepared 
from  triply  distilled  conductivity  water  which  gave  a  conductiv- 

iNoyes.  Conductivity  of  Aqueous  Solutions,  p47. 


19 


ity  of  3.9xlO-6  at  100°C-  The  values  for  A  oc  were  taken  from  the 
literature  whenever  it  was  possible.  No  corrections  were  made 
for  viscosity. 


mol  fraction 
acid 
0.0615 
0.0296 
0.0145 
0.0072 
0.0035 

0.0844 
0.0395 
0.0191 
0.0094 
0.0059 


Oxalic  Acid. 

(COOH)2 

(liters) 

Specific 

conductivity 

A 

0.328 

0.06920 

23.7 

0.656 

0.05915 

38.9 

1.312 

0.05564 

73.0 

2.624 

0.04005 

105.1 

5.248 

0.02854 

150.8 

Phosphoric  Acid.  H3P04  A  oc 

0.203  0.1256  25.5 

0.504  0.07242  36.5 

1.008  0.04920  49.6 

2.016  0.03402  68.6 

3.028  0.02813  85.2 


oc 

0.027 
0.045 
0.085 
0.122 
0.176 

=7302 
0.035 
0.050 
0.068 
0.094 
0.116 


Citric  Acid.         (CH2COOH)2COHCOOH      Aoc=8583 


0.514 
1.028 
2.057 
4.114 

8.228 


0.01248 

0.01177 

0.01001 

0.007705 

0.005710 


5.9 
12.1 
20.6 
31.7 
47.7 


d-Tartaric  Acid.   (CHOHCOOH)2 

0.433  0.1709                    7.4 

0.866  0.1593                  13.8 

1.732  0.1149                  19.9 

3.464  0.08253                28.6 

6.928  0.05946                41.2 


0.006 
0.014 
0.024 
0.037 
0.055 

Aoc=860 
0.008 
0.016 
0.023 
0.033 
0.047 


When  the  values  for   oc   are  plotted  against  the 
eentage  of  acid  smooth  curves  are    obtained.     The 
oc   for  intermediate  concentrations  were  then  found 
curves. 

ij.  Johnston,  Jour.  Am.  Chem.  Soc.  31,  1010,   (1909). 
2Noyes,  loc.  cit.  p.  262. 
3J.  Johnston  loc.  cit. 


0.0455 
0.0202 
0.0095 
0.0046 
0.0022 

0.0570 
0.0245 
0.0114 
0.0055 
0.0027 
molar  per- 
values  for 
from  these 


20 


CoTxLv.  c  t  iv  i  ti 

r 

JT 

JZT  citric 


21 


BOILING-POINT  RESULTS 

The  following  acids  have  been  used  in  this  investigation; 

Acid  lonization  constant  at  25° 

Boric  l.TxlO-9.     * 

Succinnic  0.000068     " 

Citric  0.00082.      s 

d-Tartaric  0.00097.      4 

Phosphoric  0.0070+0.0013  (1—  oc )     • 


Oxalic  0.010. 

It  is  evident  that  we  have  here  a  continuous  series  from  a 
typical  weak  acid  to  a  typical  strong  acid.  The  ionization  con- 
stants at  25°  are  given  merely  for  comparison.  In  this  work  the 
ionization  was  determined  only  for  solutions  of  the  high  concen- 
trations investigated  and  no  measurements  were  made  on  solu- 
tions dilute  enough  to  give  an  ionization  constant.7 

In  the  following  tables  oc  represents  the  degree  of  ionization, 
y  the  calculated  mol  fraction  of  solute,  and  AT  the  boiling-point 
elevation.  lonization  in  the  cases  of  boric  and  succinic 
acids  has  not  been  considered  as  it  is  sufficiently  small  to  be  neg- 
lected. In  the  diagram  the  curves  are  shown  between  concen- 
trations of  2  and  6%  molar.  All  of  the  curves,  however,  continue 
with  perfect  regularity  to  the  origin.  The  theoretical  boiling- 
point  of  pure  water  was  obtained  by  extrapolation  from  the 
curves  as  it  cannot  be  measured  directly.8  The  same  value  for  the 
boiling-point  of  water  was  obtained  from  each  curve. 


iWalker  and  Cormack,  Jour.  Chem.  Soc.  77,  5,   (1900). 

3Rivett  and  Sidgwick,  Ibid,  97,  1677,   (1910). 

sWalden.  Zeit.  Phys.  Chem.  10,  563,  (1892). 

4Grossman  and  Kramer,  Zeit.  Anorg.  Chem.,  41,  43,   (1904). 

"Kendall,  Jour.  Chem.  Soc.,  101,  1294,  (1912). 

eOstwald,  Zeit.  Phys.  Chem.  3,  241,  (1889). 

7Kendall,  Meddel,  fran  K.  Vets.  Akad.  Nobel.  Band  2,  no.  38 
(1915),  for  reasons  why  concentrated  solutions  do  not  give  an  ioniza- 
tion constant. 

«No  record  has  been  found  in  the  literature  of  anyone  obtaining 
consistent  results  for  the  boiling-point  of  water  by  immersing  the 
thermometer  in  the  liquid.  This  fact  is  suggested  in  the  calibration 
of  a  thermometer  by  calling  100 °C  the  "steam  point."  For  a  full  dis- 
cussion on  the  boiling-point  of  water  see  Preston,  Theory  of  Heat 
(1904)  p.  360. 


22 


23 

I.  Boric  Acid 

cone,  wt-  molar  y  AT 

0.5971  0.0106  0.291 

0.9656  0.0171  0.492 

1.1342  0.0200  0.578 

1.6942  0.0295  0.878 

2.2713  0.0392  1.155 

2.2741  0.0393  1.164 

2.8627  0.0490  1.476 

3.5055  0.0593  1.786 

4.1225  0.0691  2.114 

The  curve  for  boric  acid  follows  the  theoretical  curve 
throughout  the  entire  length  investigated.  No  indication  of  hy- 
dration  is  apparent,  as  we  would  expect,  since  very  weak  acids 
are  unable  to  form  addition  compounds  with  other  weak  acids — 
water  in  this  case  acting  as  an  exceedingly  weak  acid. 

II.  Succinic  Acid. 

cone,  wt-  molar  y  AT 

0.5634  0.0100  0.243 

1.1276  0.0198  0.520 

1.1322  0.0199  0.580 

1.7186  0.0300  0.827 

2.2973  0.0397  1.148 

2.2987  0.0398  1.163 

2.8755  0.0491  1.461 

2.8798  0.0493  1.463 

3.7226  0.0623  1.815 

4.0369  0.0677  1.959 

The  curve  for  succinic  acid  is  peculiar.  At  the  lowest  con- 
centrations it  falls  below  the  theoretical.  As  the  concentration 
is  increased  its  slope  gradually  becomes  steeper  and  the  curve 
rises  more  rapidly  until  it  crosses  the  theoretical  curve  at  a  con- 
centration of  about  6%  weight  molar.  On  investigating  the 
facts  known  about  this  acid  we  find  that  the  position  of  the  curve 
is  exactly  as  we  would  expect.  Succinic  acid,  being  associated 
in  aqueous  solutions,1  would  give  a  much  lower  concentration 

iPeddle  and  Turner,  loc.  cit. 


24 

than  that  calculated  on  the  assumption  of  simple  molecules.  As 
the  concentration  is  increased  succinic  acid  would  become  more 
and  more  associated.  At  the  same  time  the  influence  which  is 
probably  due  to  hydration  would  become  more  and  more  pro- 
nounced until  it  finally  overbalances  the  association  influence  and 
the  curve  rises  more  rapidly.  The  temperature  interval  through- 
out the  whole  range  is  not  great  enough  to  have  any  appreciable 
effect  on  the  association  of  the  succinic  acid  molecules. 

III.  d-Tartaric  Acid. 

cone.  wt.  molar  oc                         y  AT 

0.5628  0.023  0.0102  0.287 

1.1256  0.016  0.0202  0.625 

1.6856  0.013  0.0297  0.942 

2.3027  0.010  0.0401  1.337 

2.8598  0.008  0.0493  1.781 

3.5027  0.006  0.0596  2.265 

The  curve  for  d-tartaric  acid  diverges  widely  from  the 
theoretical,  showing  that  the  mol  fraction  of  the  acid  must  be 
much  greater  than  that  calculated  from  the  weights  of  the  com- 
ponents taken,  or  in  other  words,  that  part  of  the  solvent  has 
probably  been  removed  by  union  with  the  acid. 

IV.  Citric  Acid. 

cone.  wt.  molar  oc                         y  AT 

0.8698  0.019  0.0157  0.486 

1.0728  0.015  0.0192  0.598 

1.2736  0.013  0.0227  0.702 

1.6667  0.011  0.0294  0.950 

1.8488  0.010  0.0325  1.080 

2.0309  0.009  0.0356  1.192 

2.5847  0.006  0.0448  1.552 

3-5597  0.004  0.0605  2.346 

The  curve  for  citric,  acid  is  slightly  more  divergent  than 
that  for  tartaric  acid.1  At  25°  citric  acid  has  a  smaller  ionization 
constant  than  tartaric  acid,  showing  that  at  this  temperature  at 
least  it  is  weaker  acid.  It  might  seem  at  first  sight  as  though  our 
theory  had  failed.  An  inspection  of  the  conductivity  curves, 

ic.  f.  Juttner,  Zeit  Phys.  Chem.  38,  112,   (1901). 


25 

however,  shows  that  in  the  more  dilute  solutions — those  suffic- 
iently dilute  to  give  an  ionization  constant — at  100°,  citric  acid 
is  more  highly  ionized,  i.  e.  stronger,  than  tartaric  acid  and 
should  therefore  combine  with  water  to  a  greater  extent  as  the 
curves  indicate-1 

V.  Phosphoric  Acid. 

y  a  y 

cone.  wt.  molar  (uncorr)  (uncorr)  (corr)        AT 

0.3807  0.0068  0.117  0.0075       0.216 

0.7006  0.0124  0.096  0.0136       0.345 

1.2226  0.0214  0.066  0.0229       0.686 

1.5618  0.0275  0.061  0.0291       0.905 

1.7340  0.0304  0.058  0.0319       1.019 

1.7980  0.0313  0.057  0.0337       1.058 

2.1344  0.0370  0.053  0.0388       1.354 

2.2249  0.0385  0.051  0.0403       1.374 

2.6222  0.0449  0.048  0.0470       1.663 

2.9419  0.0510  0.046  0.0523       1.937 

3.3646  0.0571  0.044  0.0591       2.369 

3.9556  0.0664  0.041  0.0689       2.943 

In  the  case  of  phosphoric  acid  it  was  thought  best  to  show 
the  uncorrected  curve.  The  viscosity  correction  in  this  case  is 
very  large  and  the  values  for  oc  are  at  best  only  approxima- 
tions. It  is  evident,  however,  that  the  corrected  curve  would  lie 
somewhere  between  those  for  citric  acid  and  oxalic  acid. 

VI.  Oxalic  Acid. 


cone.  wt.  molar 

oc 

y 

AT 

0.5574 

0.109 

0.0110 

0.347 

0.5865 

0.106 

0.0115 

0.364 

1.0919 

0.072 

0.0206 

0.659 

1.1504 

0.070 

0.0217 

0.730 

1.4996 

0.053 

0.0276 

0.934 

1.5927 

0.050 

0.0292 

0.982 

2.0266 

0.042 

0.0368 

1.265 

2.6563 

0.036 

0.0471 

1.704 

3.0499 

0.032 

0.0535 

2.026 

iBetween  0°  and  35°  the  ionization  of  tartaric  acid  decreases 
slightly  faster  than  the  ionization  of  citric  acid.  (Jones  and  White,  Am. 
Chem.  Jour.  44,  159,  (1910)  ). 


26 

cone.  wt.  molar  oc  y  AT 

3.3099  0.030  0.0578  2.244 

3.3460  0.029  0.0584  2.310 

3.6421  0.027  0.0631  2.491 

3.9790  0.026  0.0647  2.513 

The  curve  for  oxalic  acid  shows  the  greatest  divergency 
of  the  series.  As  oxalic  acid  is  the  strongest  acid  examined,  its 
position  is  exactly  that  required  by  our  theory  of  boiling-point 
elevation. 

VII.  lodic  Acid. 

An  attempt  was  made  to  use  iodic  acid.  This  resulted  in  a 
failure  on  account  of  the  decomposition  at  the  boiling-point  of 
the  solution.  Even  in  very  dilute  solutions,  iodine  vapor  could  be 
plainly  seen  above  the  liquid.  As  there  was  no  method  of  estim- 
ating the  amount  of  decomposition  at  the  time  the  temperature 
readings  were  taken,  these  results  had  to  be  discarded. 

SUMMARY 

The  boiling-point  curves  for  a  series  of  acids  have  been  ex- 
amined, all  measurements  having  been  made  under  a  constant 
pressure.  In  every  case  examined  the  elevation  of  the  boiling- 
point  was  found  to  be  proportional  to  the  strength  of  the  acid, 
the  stronger  the  acid  the  greater  the  elevation  of  the  boiling- 
point.  This  is  taken  as  an  indication  that  the  hydration  in  solu- 
tion increases  regularly  with  the  acidic  strength,  the  abnormal- 
ity of  the  boiling-point  elevation  being  explained  on  the  assump- 
tion that  part  of  the  solvent  has  combined  with  the  solute,  thus 
giving  a  higher  concentration  of  solute  than  that  calculated 
from  the  weights  of  the  two  components. 


VITA 

Horatio  Wales  Jr.,  was  born  in  Polo,  Illinois  on  February 
12th,  1894.  He  attended  the  public  schools  in  that  city  and  was 
graduated  from  the  High  School  in  1911.  In  1915,  he  received 
the  degree  of  A.  B.,  from  Amherst  College,  and  in  1917  was 
granted  the  M.  A.  degree  at  Columbia  University.  He  was  lab- 
oratory assistant  in  Chemistry  at  Columbia  University  from  1916 
to  1918  while  a  student  at  the  University. 


----.SSSSB'"— 


«OV  10  193; 

2Uun'57B  J 
LD 


-i    :,o, 


Gaylord  Bros. 

Makers 

Syracuse,  N.  V. 
PAT.  JAM.  21 ,1908 


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